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Higher Chemistry > Unit 1 > Flashcards

Unit 1 Flashcards

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1
Q

Define activation energy ?

A

The amount of energy required to start the reaction

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2
Q

What happens when temperature is increased.

A

At higher temperatures more particles have energy equal to or above the activation energy and so therefore there is more successful collisions

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3
Q

Enthalpy means

Enthalpy change means

A

Energy

Energy change

Symbol is (H)

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4
Q

What can catalysts be used for

A

They can lower the activation energy

Make the process more economical

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5
Q

The enthalpy change is ?

A

The enthalpy change is the start to end value difference

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6
Q

Define activated complex

A

The maximum energy point along the reaction path

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7
Q

On the graphs the shaded area represents all the molecules with energy greater than the activation energy. An increase in activation energy would cause a decrease in the shaded area

A

.

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8
Q

How can covalent bond length be measured?

Covalent radius is ?

A

X-ray diffraction

The centre to the edge of the molecule. Not the bond length

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9
Q

Atomic size . Explain

A

Decreases across the periodic table - protons added moving across so there is a greater nuclear charge, with the same amount of shielding.

Increases down. - new electron shell added each time moves down
- greater shielding from the extra layer of electrons

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10
Q

What is an ion?

A

An ion is an atom which has lost or gained an electron

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11
Q

Define ionisation energy ?

A

Ionisation energy is the energy required to completely remove 1 mole of electrons from one mole of atoms in the gaseous state

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12
Q

When subbing in from data booklet make sure use state symbols

A

.

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13
Q

Ionisation energy . Explain periodic table of This!

A

Decreases down:
- more electron shells and more shielding

  • decreased attraction between outer electrons and nucleus, less energy required to move electron

Increases along:
- more protons added going across, so greater nuclear charge

  • greater attraction between nucleus and outer electron being removed, so more energy required
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14
Q

Define electro negativity !

A

A measure of the attraction an atom ya for its bonding electrons

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15
Q

High electronnegativity :

Low electro negativity :

A

High - attracts electrons strongly

Low - attracts electrons weakly

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16
Q

Electro negativity on the periodic table. Explain.

A

Decreases down

  • more electron shells being added and more shielding
  • decreases attraction between outer electron and nucleus, so lower electro negativity

Increases
- more protons being added across so greater nuclear charge

  • greater attraction between nucleus and outer electron being removed so more energy required
17
Q

Group 0

Quick facts

A
  • monatomic
  • covalent bonds
  • IMFs

Examples : helium and argon

18
Q

Cov network

A

Carbon, silicon, boron

19
Q

Fullerene is ?

A

A form of carbon which has a covalent molecular structure

20
Q

Group 7

Quick facts

A
  • bonding: covalent
  • structure : molecular
  • they are diatomic elements

Examples: fluorine, chlorine

21
Q

Non polar covalent bonds are formed when?

Polar covalent bonds are formed when?

A

When the electronegativity values of each atom are the same (>0.4)

Formed when the electronegativity values of each atom are different

22
Q

A difference in electronegativity causes a ?

A

Permanent dipole

And the more electronegative of the two molecules has the dots on the line nearest it and has a delta negative sign beside it

23
Q

Polar shape is ?

Non polar shape is

A

Asymetrical

Symmetrical

24
Q

More electrons means the higher the LDF is

A

.

25
Q

How are permanent dipoles formed?

A

Occur when electrons are unequally shared (electronegativity difference)

26
Q

When does hydrogen bonding occur ?

A

When H is directly bonded to N, O, F

Must have line in between to be a hydrogen bond

27
Q

Facts -

H2O, HF,NH3. All have higher boiling points due to the presence of hydrogen bonds between the molecules which means that more energy is required to overcome these intermolecular forces

A

.

28
Q

Why does ice float?

A

When water molecules form ice they create a large structure with large spaces between the molecules which results in it being less dense than water and so the ice floats

29
Q

Reasons why an element has a higher boiling point than another ?

(Memorise word for word this)

A

The more electrons there are means the stronger the London dispersion force is. This makes it harder to break up the bonds, the further down the halogens, as the number of electrons increase.

Example :

Sulfur has a higher number of electrons than phosphorus meaning the London dispersion force is higher, making it harder to break intermolecular bonds. S8 has more electrons than P4, and so it has a higher London dispersion force

30
Q

The more intermolecular forces, the more viscous a substance is (thicker)

A

Viscous means how thick it is

Example : syrup

31
Q

Water is more viscous than ethanol because ?

A

Water CAN form hydrogen more hydrogens bonds

32
Q

How to know what dissolves what ?

A

Polar dissolves polar

Non polar dissolves non polar

‘Like dissolves like’

33
Q

Miscibility ?

Define miscible liquids

Define immiscible liquids

A

Miscible liquids is a liquid that can mix

Immiscible liquids can’t mix (I.e. oil and water)

34
Q

The best group for being an oxidising agent is?

The best group for being a reducing agent is?

A

group 7 (halogens). Because they are very electronegative and also have a greater attraction for bonding electrons

Group 1 because they want to gain a full outer shell (metals)

35
Q

Oxidising agents accept electrons

Reducing agents give electrons

A

.

36
Q

Balancing equations steps:?

A
  • balance everything except oxygen and hydrogen
  • balance Oxygen by adding H2O
  • balance H by adding H+
  • add electrons to same side as hydrogens to make charges equal on both sides
37
Q

How are London dispersion forces formed/caused?

A

They are caused by the movement of electrons, as electrons move round the nucleus, one side of the atom becomes positive, and the other becomes negative for a fraction of a second. This causes a temporary dipole, this then causes a neighbouring atom to form an induced dipole for a fraction of a second

38
Q

Are ionic lattices polar or non polar ?

A

Polar

Higher Chemistry (12 decks)

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